Writing Lewis Formulas: The Octet Rule
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The octet rule states that representative elements usually attain stable noble gas electron
configurations in most of their compounds.
Lewis dot formulas are based on the octet rule.
We need to distinguish between bonding (or shared) electrons and nonbonding (or
unshared or lone pairs) of electrons.
N - A = S rule
– Simple mathematical relationship to help us write Lewis dot formulas.
N = number of electrons needed to achieve a noble gas configuration.
– N usually has a value of 8 for representative elements.
– N has a value of 2 for H atoms.
A = number of electrons available in valence shells of the atoms.
– A is equal to the periodic group number for each element.
– A is equal to 8 for the noble gases.
S = number of electrons shared in bonds.
A-S = number of electrons in unshared, lone, pairs.
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Writing Lewis Dot Formulas
N ever Have a Full Octet
Always Have a Full Octet
Sometimes Have a Full Octet
Sometimes Exceed a Full Octet
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Writing Lewis Formulas: The Octet Rule
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2.
For ions we must adjust the number of electrons available, A.
a.
Add one e- to A for each negative charge.
b.
Subtract one e- from A for each positive charge.
The central atom in a molecule or polyatomic ion is determined by:
a.
The atom that requires the largest number of electrons to complete its octet goes
in the center.
b.
For two atoms in the same periodic group, the less electronegative element goes in
the center.
3. Select a reasonable skeleton
a. The least electronegative is the central atom
b. Carbon makes 2,3, or 4 bonds
c. Nitrogen makes 1(rarely), 2,3, or 4 bonds
d. Oxygen makes 1, 2(usually), or 3 bonds
e. Oxygen bonds to itself only as O2 or O3, peroxides, or superoxides
f. Ternary acids (those containing 3 elements) hydrogen bonds to the oxygen, not
the central atom, except phosphates
g. For ions or molecules with more than one central atom the most symmetrical
skeleton is used
4. Calculate N, S, and A
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Count the number of electrons brought to the party (# of element times group number)
For ions we must adjust the number of electrons available.
a.
Add one e- to A for each negative charge.
b. Subtract one e- from A for each positive charge.
Select a reasonable skeleton
a. The least electronegative is the central atom
b. See prior periodic table for number of electrons involved in bonding
a. Group I 2 electrons or 1 bond
b. Group II 4 electrons or up to 2 bonds
c. Group III Al and B, 6 or 8 electrons up to 3 or 4 bonds
d. C,N,O,F must have 8 electrons (up to 4 bonds for C, 3 for N, 2 for O, and 1 bond
for F).
e. All others must have at least 8 electrons (up to 4 bonds), but may have more.
The central atom in a molecule or polyatomic ion is determined by:
a.
For ions or molecules with more than one central atom the most symmetrical
skeleton is used
b. The atom that requires the largest number of electrons to complete its octet goes in
the center.
c.
For two atoms in the same periodic group, the less electronegative element goes in
the center.
Calculate Formal charges, adjust bonds for lowest numbers (zero preferred) and allow for
resonance structures
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Writing Lewis Formulas:
The Octet Rule
• Write Lewis dot and dash formulas for hydrogen cyanide, HCN.
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Writing Lewis Formulas:
The Octet Rule
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Write Lewis dot and dash formulas for the sulfite ion, SO32-.
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Writing Lewis Formulas:
The Octet Rule
• What kind of covalent bonds, single, double, or triple, must this ion have so
that the six shared electrons are used to attach the three O atoms to the S
atom?
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Resonance
• Write Lewis dot and dash formulas for sulfur trioxide, SO3.
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Resonance
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There are three possible structures for SO32-.
– The double bond can be placed in one of three places.
oWhen two or more Lewis formulas are necessary to show the
bonding in a molecule, we must use equivalent resonance
structures to show the molecule’s structure.
oDouble-headed arrows are used to indicate resonance formulas.
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Writing Lewis Formulas:
Limitations of the Octet Rule
Write dot and dash formulas for BBr3.
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Writing Lewis Formulas:
Limitations of the Octet Rule
• Write dot and dash formulas for AsF5.
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Stereochemistry
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Stereochemistry is the study of the three dimensional shapes of molecules.
Valence Shell Electron Pair Repulsion Theory
• Commonly designated as VSEPR
• Principal originator
– R. J. Gillespie in the 1950’s
Valence Bond Theory
• Involves the use of hybridized atomic orbitals
• Principal originator
– L. Pauling in the 1930’s & 40’s
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The same basic approach will be used in every example of molecular structure
prediction:
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Polar Molecules: The Influence of Molecular Geometry
• Molecular geometry affects molecular polarity.
– Due to the effect of the bond dipoles and how they either cancel or reinforce
each other.
A
A B A
linear molecule
nonpolar
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B
A
angular molecule
polar
Polar Molecules must meet two requirements:
1. One polar bond or one lone pair of electrons on central atom.
2. Neither bonds nor lone pairs can be symmetrically arranged that their
polarities cancel.
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VSEPR Theory
• Regions of high electron density around the central atom are arranged as far apart as
possible to minimize repulsions.
• There are five basic molecular shapes based on the number of regions of high electron
density around the central atom.
• Lone pairs of electrons (unshared pairs) require more volume than shared pairs.
– Consequently, there is an ordering of repulsions of electrons around central atom.
• Criteria for the ordering of the repulsions:
1 Lone pair to lone pair is the strongest repulsion.
2 Lone pair to bonding pair is intermediate repulsion.
3 Bonding pair to bonding pair is weakest repulsion.
• Mnemonic for repulsion strengths
lp/lp > lp/bp > bp/bp
• Lone pair to lone pair repulsion is why bond angles in water are less than 109.5o.
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VSEPR Theory
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Frequently, we will describe two geometries for each molecule.
Electronic geometry is determined by the locations of regions of
high electron density around the central atom(s).
Molecular geometry determined by the arrangement of atoms
around the central atom(s).
Electron pairs are not used in the molecular geometry
determination just the positions of the atoms in the
molecule are used.
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VSEPR Theory
• Two regions of high electron density around the central atom.
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Three regions of high electron density around the central atom.
• Four regions of high electron density around the central atom.
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VSEPR Theory
• Five regions of high electron density around the central atom.
• Six regions of high electron density around the central atom.
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VSEPR Theory
• An example of a molecule that has different electronic and molecular geometries is water H2O.
• Electronic geometry is tetrahedral.
• Molecular geometry is bent or angular.
H
H C
H
H
• An example of a molecule that has the same electronic and molecular
geometries is methane - CH4.
• Electronic and molecular geometries are tetrahedral.
H
H C
H
H
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Valence Bond (VB) Theory
Regions of High Electron
Density
Electronic Geometry
Hybridization
2
Linear
sp
3
Trigonal planar
sp2
4
Tetrahedral
sp3
5
Trigonal bipyramidal
sp3d
6
Octahedral
sp3d2
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Molecular Shapes and Bonding
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In the next sections we will use the following terminology:
A = central atom
B = bonding pairs around central atom
U = lone pairs around central atom
For example:
AB3U designates that there are 3 bonding pairs and 1 lone pair around the central
atom.
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Linear Electronic Geometry:AB2 Species (No
Lone Pairs of Electrons on A)
Be
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1s
2s


2p
1s
 
sp hybrid 2p


Trigonal Planar Electronic Geometry: AB3
Species (No Lone Pairs of Electrons on A)
B
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1s 2s 2p
  
1s
 
sp2 hybrid
  
Tetrahedral Electronic Geometry: AB4
Species (No Lone Pairs of Electrons on A)
2s
C [He] 
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2p

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Tetrahedral Electronic Geometry: AB4 Species
Valence Bond Theory (Hybridization)
C [He]
2s

2p
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four sp3 hybrids

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Tetrahedral Electronic Geometry: AB3U Species
2s
N [He]

2p four sp3 hybrids

Tetrahedral Electronic Geometry: AB2U2 Species
O [He]
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2s
 
3
2p four sp hybrids
 
Tetrahedral Electronic Geometry: ABU3 Species (Three
Lone Pairs of Electrons on A)
Valence Bond Theory (Hybridization)
F [He]
four sp3 hybrids
   
2s
2p
  
··
H
F
··
··
tetrahedral
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Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2,
and AB2U3
4s
4p
   
4d
As [Ar] 3d10
_______________

five sp3 d hybrids
4d
    
____________
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Compounds Containing
Double Bonds
Valence Bond Theory (Hybridization)
C atom has four electrons.
Three electrons from each C atom are in sp2 hybrids.
One electron in each C atom remains in an unhybridized p orbital
2s 2p
three sp2 hybrids 2p
C    


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An sp2 hybridized C atom has this shape.
Remember there will be one electron in each of the three lobes.
Top view of
an sp2 hybrid
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Compounds Containing
Double Bonds
• The single 2p orbital is perpendicular to the trigonal planar sp2 lobes.
The fourth electron is in the p orbital.
Side view of sp2 hybrid
with p orbital included.
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Compounds Containing
Double Bonds
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Two sp2 hybridized C atoms plus p orbitals in proper orientation to form C=C
double bond.
• The portion of the double
bond formed from the
head-on overlap of the sp2
hybrids is designated as a
s bond.
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The other portion of the
double bond, resulting from the
side-on overlap of the p
orbitals, is designated as a p
bond.
Compounds Containing Triple Bonds
A s bond results from the head-on overlap of two sp hybrid
orbitals.
The unhybridized p orbitals form two p bonds.
Note that a triple bond consists of one s and two p bonds.
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Summary of Electronic & Molecular
Geometries
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